Arrange the following substances in the expected order of increasing melting point: \(\mathrm{KI}\), \(\mathrm{Ne}, \mathrm{K}_{2} \mathrm{SO}_{4}\) \(\mathrm{C}_{3} \mathrm{H}_{8}, \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}, \mathrm{MgO}, \mathrm{CH}_{2} \mathrm{OHCHOHCH}_{2} \mathrm{OH}\)

Understanding intermolecular forces is crucial when predicting the melting point order of different substances. These forces are the interactions that occur between molecules and include various types of bonds like London dispersion forces, dipole-dipole interactions, and hydrogen bonds. Let's break them down:

London Dispersion Forces: These are the weakest intermolecular forces caused by the random movement of electrons that create temporary dipoles. Substances with London dispersion forces usually have lower melting points, such as noble gases and nonpolar molecules.
Dipole-Dipole Interactions: These occur when molecules with permanent dipoles interact. They are stronger than dispersion forces and are typical in polar molecules.
Hydrogen Bonding: This specific type of intermolecular force arises when a hydrogen atom bonded to a highly electronegative atom like oxygen, nitrogen, or fluorine, is electrostatically attracted to another electronegative atom. This type of bond is particularly strong compared to other intermolecular forces and greatly influences the melting point.
In our exercise, Ne would have the lowest melting point due to weak dispersion forces, followed by C_{3}H_{8}, a nonpolar molecule also with dispersion forces. Substances that have hydrogen bonding, such as alcohols and sugars, in this case, CH_{3}CH_{2}OH and CH_{2}OHCHOHCH_{2}OH, would have higher melting points because hydrogen bonds are stronger compared to typical dispersion forces.

Ionic compounds are formed by the electrostatic attraction between oppositely charged ions, usually a metal and a non-metal. These compounds have high melting and boiling points due to the strength of the ionic bonds that hold the ions together in a lattice structure. The stronger the charge of the ions and the smaller their sizes, the stronger the ionic bonds become.

For example, in our exercise, KI, K_{2}SO_{4}, and MgO are ionic compounds. Among them, MgO has the highest melting point because magnesium and oxygen ions have a larger charge than the ions in KI and K_{2}SO_{4}. This powerful attraction requires a significant amount of energy to overcome, explaining why ionic compounds tend to have much higher melting points than substances dominated by other types of intermolecular forces.

Let's delve deeper into hydrogen bonding to appreciate why it profoundly affects the melting points of certain compounds. When hydrogen is attached to atoms like oxygen, it can partially expose its nucleus because of the small size of hydrogen and the significant electronegativity difference between hydrogen and the adjacent atom. This leads to a strong dipole and the ability for the hydrogen to act like a magnetic 'hook' that attaches to other electronegative atoms nearby.

Alcohols and sugars, such as CH_{3}CH_{2}OH (ethanol) and CH_{2}OHCHOHCH_{2}OH (a triol sugar), show this property. Hydrogen bonding makes these substances stick together more tightly and therefore increases their melting points compared to other organic compounds without hydrogen bonds like the alkanes (e.g., C_{3}H_{8}, propane) in our exercise. The presence of multiple -OH groups in a sugar molecule like CH_{2}OHCHOHCH_{2}OH contributes to an even higher melting point, as there are more spots for these strong intermolecular attractions to occur.